Chemistry for Kids: Elements - The Noble Gases Elements for Kids Noble Gases The noble gases are a group of elements in the periodic table. They are located to the far right of the periodic table and make up the eighteenth column. Elements in the noble gas family have atoms with a full outer shell of electrons. They are also called the inert gases. What elements are noble gases? The elements that make up the family of noble gases include helium , neon , argon , krypton, xenon, and radon. What are the similar properties of noble gases? Noble gases share many similar properties including: A full outer shell of electrons. Helium has two electrons in its outer shell and the rest have eight electrons. Because of their full outer shells, they are very inert and stable. This means they don't tend to react with other elements to form compounds. They are gases under standard conditions. They are colorless and odorless. Their melting and boiling points are close together giving them a very narrow liquid range. Abundance Helium is the second most abundant element in the universe after hydrogen. Helium makes up about 24% of the mass of the elements in the universe.

Neon is the fifth most abundant and argon is the eleventh. On Earth, the noble gases are fairly rare with the exception of argon. Argon makes up just under 1% of the Earth's atmosphere, making it the third most abundant gas in the atmosphere after nitrogen and oxygen. Interesting Facts about Noble Gases Because helium is non-flammable it is much safer to use in balloons than hydrogen. Krypton gets its name from the Greek word "kryptos" meaning "the hidden one." Many of the noble gases were either discovered or isolated by Scottish chemist Sir William Ramsay. Helium has the lowest melting and boiling points of any substance. All of the noble gases except for radon have stable isotopes. Neon signs do not use just neon gas, but a mixture of different noble gases and other elements to create bright lights of different colors. Noble gases are often used to create a safe or inert atmosphere due to their stable nature. Xenon gets its name from the Greek word "xenos" which means "stranger or foreigner." More on the Elements and the Periodic Table Noble Gases | IAS Noble Gases Tweet on Twitter The noble gases are a group of elements in the periodic table .

They are located to the far right of the periodic table and make up the eighteenth column. Elements in the noble gas family have atoms with a full outer shell of electrons. They are also called the inert gases. The elements that make up the family of noble gases include helium, neon, argon, krypton, xenon, and radon. What are the similar properties of noble gases? Noble gases share many similar properties including: A full outer shell of electrons. Helium has two electrons in its outer shell and the rest have eight electrons. Because of their full outer shells, they are very inert and stable. This means they don’t tend to react with other elements to form compounds. They are gases under standard conditions. They are colourless and odourless. Their melting and boiling points are close together giving them a very narrow liquid range. Abundance Helium is the second most abundant element in the universe after hydrogen. Helium makes up about 24% of the mass of the elements in the universe . Neon is the fifth most abundant and argon is the eleventh.On Earth , the noble gases are fairly rare with the exception of argon.

Argon makes up just under 1% of the Earth’s atmosphere, making it the third most abundant gas in the atmosphere after nitrogen and oxygen. Also Read: Halogens Radon Radon is the heaviest of the noble gases and was discovered due to the decay of radium. Uranium and radium are naturally present in rock and soil. Radon can get concentrated in unventilated indoor areas built into the soil such as basements. Due to its colourless property, it can’t be detected easily. Radon accumulating indoors is a serious health hazard. This is because when radon breaks down, it produces charged particles that become a part of the dust, and are dangerous when inhaled. This hazard is also posed in structures constructed from certain granites. Radon does have its uses too. Radon-222, the most common isotope, is used in the treatment of malignant tumours. Interesting Facts about Noble Gases Because helium is non-flammable it is much safer to use in balloons than hydrogen. Krypton gets its name from the Greek word “kryptos” meaning “the hidden one.” Many of the noble gases were either discovered or isolated by Scottish chemist Sir William Ramsay.

Helium has the lowest melting and boiling points of any substance. All of the noble gases except for radon have stable isotopes. Neon signs do not use just neon gas, but a mixture of different noble gases and other elements to create bright lights of different colours. Noble gases are often used to create a safe or inert atmosphere due to their stable nature. Xenon gets its name from the Greek word “xenos” which means “stranger or foreigner.” Sir William Ramsay (1852-1916) was a British chemist, and was the first to successfully isolate all of the noble gases, except radon. He also discovered 4 of them, namely neon, argon, krypton, and xenon. Pierre Janssen discovered helium in 1868 during a solar eclipse in the spectrum of the corona of the sun . However, Ramsay was the first to isolate it. Ramsay and the British physicist Lord Raleigh found that nitrogen removed from the atmosphere was denser than nitrogen formed by the decomposition of ammonium nitrate. This led to the discovery of argon in 1894. Helium was discovered by Janssen in 1895.

Through the attempts of Ramsay and his assistant Morris Travers, krypton, neon and xenon were found from the atmosphere by fractional distillation. Radon was discovered by a German chemist, Friedrich Ernst Dorn in 1900. List of Noble Gases | eHow List of Noble Gases neon sign image by pearlguy from Fotolia.com The six noble gases are all found in group 18 of the periodic table. Noble gases are different from other elements because they have a full set of electrons in their outer shell. Most gases have too few or too many. This complete electron shell makes noble gases very stable, a fact that once saw them referred to as "inert gases." However, some of the noble gases are now known to take part in chemical reactions. Helium Helium has an atomic number of 2 and is represented by the symbol He. The gas was discovered in 1895, but, according to the BBC website, its existence had been previously proposed by Pierre Jules Cesar and Sir Joseph Norman Lockyer. The two men that named the gas "helium" after the Greek word for sun. Helium is often used in balloons due to its lighter than air properties.

It is also used in deep sea diving equipment to dilute a diver's oxygen supply. Neon Neon has an atomic number of 10 and the symbol Ne. It was discovered in 1898 by two British chemists, William Ramsay and Morris W. Travers, and named after the Greek word for new, "neos." Neon is often used in lighting and fluorescent signs. The gas itself is colorless but glows in an electric discharge. Argon Argon has an atomic number of 18 and the symbol Ar. It was isolated in 1894 by Lord Rayleigh and William Ramsay, becoming the first of the noble gases to be discovered. The stable nature of the gas earned it the name "argon," the Greek word for inert. According to the Royal Society of Chemistry website, "Argon is widely used to provide an inert atmosphere for high-temperature metallurgical processes." Like neon, it is also used in lighting. Krypton Krypton is represented by the symbol Kr and has an atomic number of 36. Krypton was discovered in 1898 by William Ramsay and Morris W. Travers, just a few weeks before they discovered neon.

It was named after the Greek word "kryptos," meaning hidden. The gas is used in fluorescent lighting and in various types of photographic flash bulbs. Xenon Xenon has an atomic number of 54 and the symbol Xe. It was discovered by Ramsey and Travers in 1898, shortly after their discovery of neon and krypton, and named after the Greek word for stranger. Xenon is found naturally in the atmosphere, but only in minute quantities. According to the PeriodicTable.org website, xenon is used in photographic flashes, general anesthetic and stroboscopic, bactericidal, and laser-pumping lamps. Radon Radon, the last noble gas to be isolated, has an atomic number of 86 and the symbol Rn. It was initially observed in the 1890s before being isolated by Friedrich Ernst Dorn in 1900. It was properly classified in 1908 by William Ramsay. Radon is colorless, odorless, tasteless and, unlike the other noble gases, radioactive. It occurs naturally from decaying radium and is considered hazardous due to its radioactivity. It is used in radiotherapy for the treatment of cancer.

Noble gases facts, information, pictures | Encyclopedia.com articles about Noble gases Isolation of the Noble Gases HELIUM. Helium is an unusual element in many respects—not least because it is the only element to have first been identified in the Solar System before it was discovered on Earth. This is significant, because the elements on Earth are the same as those found in space: thus, it is more than just an attempt at sounding poetic when scientists say that humans, as well as the world around them, are made from "the stuff of stars." In 1868, a French astronomer named Pierre Janssen (1824-1907) was in India to observe a total solar eclipse. To aid him in his observations, he used a spectroscope, an instrument for analyzing the spectrum of light emitted by an object. What Janssen's spectroscope showed was surprising: a yellow line in the spectrum, never seen before, which seemed to indicate the presence of a previously undiscovered element. Janssen called it "helium" after the Greek god Helios, or Apollo, whom the ancients associated with the Sun .

Janssen shared his findings with English astronomer Sir Joseph Lockyer (1836-1920), who had a worldwide reputation for his work in analyzing light waves. Lockyer, too, believed that what Janssen had seen was a new element, and a few months later, he observed the same unusual spectral lines. At that time, the spectroscope was still a new invention, and many members of the worldwide scientific community doubted its usefulness, and therefore, in spite of Lockyer's reputation, they questioned the existence of this "new" element. Yet during their lifetimes, Janssen and Lockyer were proven correct. NEON, ARGON, KRYPTON, AND XENON. They had to wait a quarter century, however. In 1893, English chemist Sir William Ramsay (1852-1916) became intrigued by the presence of a mysterious gas bubble left over when nitrogen from the atmosphere was combined with oxygen. This was a phenomenon that had also been noted by English physicist Henry Cavendish (1731-1810) more than a century before, but Cavendish could offer no explanation.

Ramsay, on the other hand, had the benefit of observations made by English physicist John William Strutt, Lord Rayleigh (1842-1919). Up to that time, scientists believed that air consisted only of oxygen, carbon dioxide, and water vapor. However, Rayleigh had noticed that when nitrogen was extracted from air after a process of removing those other components, it had a slightly higher density than nitrogen prepared from a chemical reaction. In light of his own observations, Ramsay concluded that whereas nitrogen obtained from chemical reactions was pure, the nitrogen extracted from air contained trace amounts of an unknown gas. Ramsay was wrong in only one respect: hidden with the nitrogen was not one gas, but five. In order to isolate these gases, Ramsay and Rayleigh subjected air to a combination of high pressure and low temperature, allowing the various gases to boil off at different temperatures. One of the gases was helium—the first confirmation that the element existed on Earth—but the other four gases were previously unknown. The Greek roots of the names given to the four gases reflected scientists' wonder at discovering these hard-to-find elements: neos (new), argos (in active), kryptos (hidden), and xenon (stranger).

RADON. Inspired by the studies of Polish-French physicist and chemist Marie Curie (1867-1934) regarding the element radium and the phenomenon of radioactivity (she discovered the element, and coined the latter term), German physicist Friedrich Dorn (1848-1916) became fascinated with radium. Studying the element, he discovered that it emitted a radioactive gas, which he dubbed "radium emanation." Eventually, however, he realized that what was being produced was a new element. This was the first clear proof that one element could become another through the process of radioactive decay. Ramsay, who along with Rayleigh had received the Nobel Prize in 1904 for his work on the noble gases, was able to map the new element's spectral lines and determine its density and atomic mass. A few years later, in 1918, another scientist named C. Schmidt gave it the name "radon." Due to its behavior and the configuration of its electrons, chemists classified radon among what they continued to call the "inert gases" for another half-century—until Bartlett's preparation of xenon compounds in 1962.

Presence of the Rare Gases on Earth IN THE ATMOSPHERE. Though the rare gases are found in minerals and meteorites on Earth, their greatest presence is in the planet's atmosphere. It is believed that they were released into the air long ago as a by-product of decay on the part of radioactive materials in the Earth's crust. Within the atmosphere, argon is the most "abundant"—in comparative terms, given the fact that the "rare gases" are, by definition, rare. Nitrogen makes up about 78% of Earth's atmosphere and oxygen 21%, meaning that these two elements constitute fully 99% of the air above the Earth. Argon ranks a distant third, with 0.93%. The remaining 0.07% is made up on water vapor, carbon dioxide, ozone (O3), and traces of the noble gases. These are present in such small quantities that the figures for them are not typically presented as percentages, but rather in terms of parts per million (ppm). The concentrations of neon, helium, krypton, and xenon in the atmosphere are 18, 5, 1, and 0.09 ppm respectively. IN THE SOIL.

Radon in the atmosphere is virtually negligible, which is a fortunate thing, in light of its radioactive qualities. Few Americans, in fact, even knew of its existence until 1988, when the United States Environmental Protection Agency (EPA) released a report estimating that some ten million American homes had potentially harmful radon levels. This set off a scare, and during the late 1980s and 1990s, sales of home radon detectors boomed. Meanwhile, the federal government increased concerns with additional reports, advising people to seal their basements and ventilate their homes if radon exceeded certain levels. A number of scientists have disputed the government's claims, yet some regions of the United States appear to be at relatively high risk due to the presence of radon in the soil. The element seems to be most plentiful in soils containing high concentrations of uranium. If radon is present in a home that has been weather-sealed to improve the efficiency of heating and cooling systems, it is indeed potentially dangerous to the residents. Chinese scientists in the 1960s made an interesting discovery regarding radon and its application to seismography, or the area of the earth sciences devoted to studying and predicting earthquakes.

Radon levels in groundwater, the Chinese reports showed, rise considerably just before an earthquake. Since then, the Chinese have monitored radon concentrations in water, and used this data to predict earthquakes. EXTRACTING RARE GASES. Radon, in fact, is not the only rare gas that can be obtained as the result of radioactive decay: in 1903, Ramsay and British chemist Frederick Soddy (1877-1956) showed that the breakdown of either uranium or radium results in the production of helium atoms (beta particles). A few years later, English physicist Ernest Rutherford (1871-1937) demonstrated that radiation carrying a positive electrical charge (alpha rays) was actually a stream of helium atoms stripped of an electron. Many of the noble gases are extracted by liquefying air—that is, by reducing it to temperatures at which it assumes the properties of a liquid rather than a gas. By controlling temperatures in the liquefied air, it is possible to reach the boiling point for a particular noble gas and thereby extract it, much as was done when these gases were first isolated in the 1890s. THE UNIQUE SITUATION OF HELIUM.

Helium is remarkable, in that it only liquefies at a temperature of −457.6°F (−272°C), just above absolute zero. Absolute zero is the temperature at which the motion of atoms or molecules comes to a virtual stop, but the motion of helium atoms never completely ceases. In order to liquefy it, in fact, even at those low temperatures, it must be subjected to pressures many times that exerted by Earth's atmosphere. Given these facts, it is difficult to extract helium from air. More often, it is obtained from natural gas wells, where it is present in relatively large concentrations—between 1% and 7% of the natural gas. The majority of the Earth's helium supply belongs to the United States, where the greatest abundance of helium-supplying wells are in Texas , Oklahoma , and Kansas . During World War II , the United States took advantage of this supply of relatively inexpensive helium to provide buoyancy for a fleet of airships used for reconnaissance. There is one place with an abundant supply of helium, but there are no plans for a mining expedition any time soon. That place is the Sun, where the nuclear fusion of hydrogen atoms creates helium.

Indeed, helium seems to be the most plentiful element of all, after hydrogen, constituting 23% of the total mass of the universe. Why, then, is it so difficult to obtain on Earth? Most likely because it is so light in comparison to air; it simply floats off into space. Applications for the Noble Gases RADON, ARGON, KRYPTON, AND XENON. Though radon is known primarily for the hazards it poses to human life and well-being, it has useful applications. As noted above, its presence in groundwater appears to provide a possible means of predicting earthquakes. In addition, it is used for detecting leaks, measuring flow rates, and inspecting metal welds. One interesting use of argon and, in particular, the stable isotope argon-40, is in dating techniques used by geologists, paleontologists, and other scientists studying the distant past. When volcanic rocks are subjected to extremely high temperatures, they release argon, and as the rocks cool, argon-40 accumulates.

Because argon-40 is formed by the radioactive decay of a potassium isotope, potassium-40, the amount of argon-40 that forms is proportional to the rate of decay for potassium-40. The latter has a half-life of 1.3 billion years, meaning that it takes 1.3 billion years for half the potassium-40 originally present to be converted to argon-40. Using argon-40, paleontologists have been able to estimate the age of volcanic layers above and below fossil and artifact remains in east Africa. Krypton has a number of specialized applications—for instance, it is mixed with argon and used in the manufacture of windows with a high level of thermal efficiency. Used in lasers, it is often mixed with a halogen such as fluorine. In addition, it is also sometimes used in halogen sealed-beam headlights. Many fans of Superman, no doubt, were disappointed at some point in their lives to discover that there is no such thing as "kryptonite," the fictional element that caused the Man of Steel to lose his legendary strength. Yet krypton—the real thing—has applications that are literally out of this world.

In the development of fuel for space exploration, krypton is in competition with its sister element, xenon. Xenon offers better performance, but costs about ten times more to produce; thus krypton has become more attractive as a fuel for space flight. In addition to its potential as a space fuel, xenon is used in arc lamps for motion-picture film projection, in high-pressure ultraviolet radiation lamps, and in specialized flashbulbs used by photographers. One particular isotope of xenon is utilized for tracing the movement of sands along a coastline. Xenon is also applied in high-energy physics for detecting nuclear radiation in bubble chambers. Furthermore, neuroscientists are experimenting with the use of xenon in diagnostic procedures to clarify x-ray images of the human brain. NEON. Neon, of course, is best-known for its application in neon signs, which produce an eye-catching glow when lit up at night. French chemist Georges Claude (1870-1960), intrigued by Ramsay's discovery of neon, conducted experiments that led to the development of the neon light in 1910. That first neon light was simply a glass tube filled with neon gas, which glowed a bright red when charged with electricity.

Claude eventually discovered that mixing other gases with neon produced different colors of light. He also experimented with variations in the shapes of glass tubes to create letters and pictures. By the 1920s, neon light had come into vogue, and it is still popular today. Modern neon lamps are typically made of plastic rather than glass, and the range of colors is much greater than in Claude's day: not only the gas filling, but the coating inside the tube, is varied, resulting in a variety of colors from across the spectrum. Though the neon sign is its best-known application, neon is used for many other things. Neon glow lamps are often used to indicate on/off settings on electronic instrument panels, and lightweight neon lamps are found on machines ranging from computers to voltage regulators. In fact, the first practical color television, produced in 1928, used a neon tube to produce the red color in the receiver. Green came from mercury, but the blue light in that early color TV came from another noble gas, helium. HELIUM. Helium, of course, is widely known for its use in balloons—both for large airships and for the balloons that have provided joy and fun to many a small child.

Though helium is much more expensive than hydrogen as a means of providing buoyancy to airships, hydrogen is extremely flammable, and after the infamous explosion of the airship Hindenburg in 1937, helium became the preferred medium for airships. As noted earlier, the United States military made extensive use of helium-filled airships during the World War II. The use of helium for buoyancy is one of the most prominent applications of this noble gas, but far from the only one. In fact, not only have people used helium to go up in balloons, but divers use helium for going down beneath the surface of the ocean. In that situation, of course, helium is not used for providing buoyancy, but as a means of protecting against the diving-related condition known as "the bends," which occurs when nitrogen in the blood bubbles as the diver rises to the surface. Helium is mixed with oxygen in diver's air tanks because it does not dissolve in the blood as easily as nitrogen. Among the most fascinating applications of helium relate to its extraordinarily low freezing point.

Helium has played a significant role in the low-temperature science known as cryogenics, and has found application in research concerning superconductivity: the use of very low temperatures to develop materials that conduct electrical power with vastly greater efficiency than ordinary conductors. Close to absolute zero, helium transforms into a highly unusual liquid unlike any known substance, in that it has no measurable resistance to flow. This means that it could carry an electrical current hundreds of times more efficiently than a copper wire. WHERE TO LEARN MORE "The Chemistry of the Rare Gases" (Web site). < (May 13, 2001). "Homework: Science: Chemistry: Gases" Channelone.com (Web site). < (May 12, 2001). Knapp, Brian J.; David Woodroffe; David A. Hardy. Elements. Danbury, CT: Grolier Educational, 2000. Mebane, Robert C. and Thomas R. Rybolt. Air and Other Gases. Illustrations by Anni Matsick. New York : Twenty-First Century Books, 1995. "Noble Gases" Xrefer.com (Web site).

< (May 13, 2001). Rare Gases. Praxair (Web site). < (May 13, 2001). Stwertka, Albert. Superconductors: The Irresistible Future. New York: F. Watts, 1991. Taylor, Ron. Facts on Radon and Asbestos. Illustrated by Ian Moores. New York: F. Watts, 1990. Cite this article Pick a style below, and copy the text for your bibliography. MLA C&EN: IT'S ELEMENTAL: THE PERIODIC TABLE - THE NOBLE GASES THE NOBLE GASES NEIL BARTLETT , UNIVERSITY OF CALIFORNIA, BERKELEY The Periodic Table of the Elements, as set out by Dmitry Mendeleyev or Lother Meyer, had not allowed for the discovery of a group of elements between the highly electronegative halogens and the electropositive alkalis. So the discovery of the unreactive monatomic gas argon by Lord Rayleigh and William Ramsay in 1895 came as a total surprise.

James Dewar even indicated that the new gas could be an allotrope of nitrogen, N3, a suggestion seconded by Mendeleyev! Within three years, however, Ramsay (and coworker Morris Travers) had also discovered helium, neon, krypton, and xenon. They established their monatomic and unreactive nature. With the advent of the Rutherford-Bohr atom (1913), electron configurations of these unreactive monatomic elements soon came to have a central role in the emerging electronic theories of chemical bonding. In their 1916 papers, both G. N. Lewis and W. Kossel pointed to the electron configurations of these elements as especially stable. In each theory, the chemical properties of atoms of other elements were tied to the gain or loss of electrons from the configuration of the nearest monatomic gas. So successful were these theories in accounting for a wide range of chemical properties of the elements that the monatomic-gas electron configurations came to be thought of as chemically inviolate. This was fostered by early failures to make compounds of the gases (including an attempt by Henri Moissan to prepare an argon fluoride in 1895).

Nevertheless, in his classic 1916 paper, Kossel made an astute observation relevant to the chemical reactivity of these elements. On the basis of the first ionization potentials of the gases, Kossel noted that xenon was most likely to have the capability of forming fluorides and oxides. He also allowed that a krypton fluoride might be made. Similar predictions were made later, by Andreas von Antropoff (1924) and by Linus C. Pauling (1932), based on chemical trends in the periodic table. These predictions led D. M. Yost (with student A. L. Kaye) to attempt (in 1933) a xenon fluoride synthesis. That attempt failed. So matters rested until 1962. It was the discovery of the remarkable oxidizing properties of platinum hexafluoride in making the salt O2+PtF6 that led (via the recognition that O2 and Xe have nearly the same first ionization potentials) to the oxidation of xenon by PtF6. Later in 1962, Howard H. Claasen, Henry Selig, and John G.

Malm, at Argonne National Laboratory prepared XeF4. Syntheses of XeF2, XeF6, XeOF4, XeO2F2, XeO3, and perxenates (XeO64 salts) were quickly reported from there and elsewhere. Even the highly unstable tetrahedral tetroxide XeO4 was made (J. L. Houston, 1964). A fluoride of krypton, prepared and correctly identified as KrF2 (George C. Pimentel and J. J. Turner, 1963), was first reported as KrF4 (Aristid V. Grosse and coworkers, 1963), but no compound above Kr(II) has ever been established. Although the easier ionization of radon leads one to expect the most extensive chemistry for that element, the high instability of even the most stable isotope has severely limited studies of it. L.

Stein, of Argonne, established (in 1962) the existence of a fluoride--probably RnF2--but he and others were unable to confirm the existence of oxides or relatives of the perxenates. The first ionization potentials of the noble gases provide a measure of how firmly the outer electrons are held by the effective nuclear charge. This hardness or softness of the valence electron set correlates well with the physical properties of the gases. The first ionization potentials are as follows: He, 24.6; Ne, 21.6; Ar, 15.8; Kr, 14.0; Xe, 12.1; and Rn, 10.7 eV. Clearly, helium has the least polarizable electron cloud. This accounts for its low melting and boiling points and its low solubility in aqueous media, and hence its application in mitigating the bends as a diluent for oxygen in deep-water diving. In contrast, highly polarizable xenon has high solubility and is an excellent anesthetic. The low ionization potentials of the heavier gases also account for their chemistry.

In all of the known chemical compounds of the noble gases, the noble-gas atom has a net positive charge. We can take the difluorides as representative. In each, the noble-gas atom can be viewed as having lost an electron to the two F ligands. One way to picture this is to make the pseudo halogen Ng+ (transferring the electron to an F to make F), the Ng+ combining with the second F to make the classical electron-pair-bound [NgF]+ (Ng = noble gas). Each difluoride is then represented as a resonance hybrid of the canonical forms {F[NgF]+} and {[FNg]+F}. Bond energies in the isoelectronic relatives of the NgF+ species, ClF, BrF, and IF, do not vary greatly, so it is reasonable to assume the same to hold for the NgF+ species. So, in making the difluoride from the constituent atoms via such species, the energy term that changes most from one Ng to another is the ionization potential.

In support of this, we note that the heats of atomization (in kilocalories per mole) of XeF2 (65) and KrF2 (23) differ by almost the same energy as the first ionization potentials, about 44 kcal per mol. On this basis, since the first ionization potential of Ar is 42 kcal per mol higher than that of Kr, it can be supposed that the heat of atomization of ArF2 would be less than that of KrF2 by roughly that amount. Since the heat of atomization of KrF2 is only 23 kcal per mol, this implies that ArF2 cannot be made. Because of the high electronegativity of the positively charged noble-gas centers, any ligands for those centers must themselves be electronegative. The ligands must also be small to provide high Coulomb energy. The known chemistry is in accord with this. For compounds that can be prepared and manipulated at ordinary laboratory temperatures, ligands are, so far, few. They are, for xenon: oxygen (S. M. Williamson et al., 1963; N.

Bartlett, and coworkers, 1969), nitrogen (D. D. DesMarteau, 1981) and carbon (D. Naumann; H. J. Frohn, 1989); and for krypton (G. J. Schrobilgen, 198889): oxygen, and nitrogen. In all, the nitrogen and carbon ligands, to be effective, require linkage to other electronegative centers. The large size of the chlorine ligand means that XeCl2 can only be made and used at cryogenic temperatures. The high bond energy of O2 also leads to high thermodynamic instability of all oxides. XeO is only bound with respect to singlet oxygen, most likely 1D [O]. (E. H. Appelman of Argonne used XeF2 in aqueous solution to prepare the first examples of perbromates, and this could have involved attack of bromate by XeO as a singlet oxygen carrier.) Because of its lower effective nuclear charge, xenon is more easily oxidized than krypton. It is also a better Lewis base.

Krypton is not oxidized by PtF6, and in the solvent anhydrous HF (aHF), forms no complexes with transition-metal ions as xenon does. The HF-solvated Ag2+ ion (silver has the highest second ionization potential of any transition metal) oxidizes Xe at ordinary temperatures (Zemva et al., 1990) to make XeF2. With the solvated Au2+ ion, however (Konrad Seppelt, 2001 03), no oxidation of Xe occurs. Xenon is a sufficiently good Lewis base, however, to coordinate with Au2+ to form a variety of coordination complexes. It similarly complexes with Au3+ and with other metal ions. It is possible that Kr and Ar could also act as Lewis bases, but evidently they cannot compete with HF. Many recent findings, including the first evidence for an argon compound, have come from matrix-isolation studies at the University of Helsinki in Finland (Markku Räsänen and coworkers). These studies have established the existence of a large variety of novel compounds, all stable up to 40 K.

Included are HXeOH, HXeCCH, HKrCN, HKrCCH, and HArF. The last requires comment, because of the nonexistence of ArF2. In all of these compounds, the vibrational spectroscopic findings indicate that the canonical form ([HNg]+Y) contributes importantly to the binding of the molecules. The tiny proton is highly electronegative, and it bonds covalently to Ng in these molecules. The proton affinities of the noble gases are the following: He, 1.8; Ne, 2.2; Ar, 3.0; Kr, 4; and Xe, 6 eV. Attachment of a proton to the more polarizable gases therefore gives significant energy toward bonding. Of course, this requires a small electronegative coligand, of which F has no superior; but OH, CN, and CCH are also small ligands with high electron affinities. All of the noble-gas compounds are easily reduced, and the compounds can effectively carry their ligands as radical species, readily available to more reactive elements.

For example, KrF2, which is less bound than F2 itself, will oxidize Xe to XeF6! More electronegative cationic species such as KrF+ are even more potent (for example, the synthesis of BrF6+, Ronald J. Gillespie, and Schrobilgen, 1974). Some of the potential reagents are very fragile and will need to be used at cryogenic temperatures, but the elimination of an atom of an almost inert gas can simplify the chemistry. COURTESY OF NEIL BARTLETT START OF SOMETHING NEW Bartlett's oxidation of xenon by PtF6, shown here, sparked the field of noble-gas chemistry. Neil Bartlett is an emeritus professor in the department of chemistry, University of California, Berkeley, and an emeritus senior scientist in the Chemical Sciences Division of Lawrence Berkeley National Laboratory. Bartlett succeeded in preparing the first noble-gas compound in 1962.The Royal Society (London) awarded him its Davy Medal in 2002.

Clean Tech Funder: About Noble Gases Clean Tech Funder Noble Gases Properties The noble gases are the elements in the right-hand column of the periodic table . They are called "noble" for being chemically inert. None of the noble gases react with any other element except in extreme artificial conditions. Helium is a very light gas (only hydrogen is lighter), and the gases become progressively more dense down the table. They are gaseous to low or very low temperatures as well. Xenon condenses at 165K (-163 °F). Helium condenses at 4K (-452 °F). That is 4 degrees above absolute zero, the lowest boiling point of any element. Noble gases are used in place of vacuuum when a nonreactive atmosphere is required. For example, argon may be played over a welding site, and xenon may be used to fill an incandescent bulb. These gases fluoresce in various colors when ionized by a strong electric current. A neon tube glows orange-red, other gases glow with different colors, and the glass itself can be tinted for even greater variety.

Helium Helium is the second-lightest element and the first to be discovered off the planet. The development of spectroscopy around 1850 led to the discovery of such elements as caesium from the light that they emit or absorb. In 1868, the French astronomer Pierre-Jules-César Janssen observed a total solar eclipse in India. His spectroscope showed an unfamiliar line in the light spectrum. English astronomer Norman Lockyer identified it as a new element, which he named helium after the greek god Helios. This discovery was controversial until helium was detected in various minerals on earth around 1890. Helium is a major component of stars. On earth, it occurs in minerals not because it is bound chemically but because the helium nucleus is a product of radioactive decay of much larger atoms. This helium collects underground in pockets of natural gas and it is vented into the atmosphere. Helium is used to fill balloons and blimps. It has unique properties at temperatures approaching absolute zero. Helium tanks are used to cool anything else to near that temperature, and helium is the subject of many cryogenic experiments.

Neon Neon was discovered after krypton (see below) by Ramsay and Travers. They separated trace components from atmospheric argon, one of which produced a crimson spectrum. Neon means "new". Neon was the first gas used in so-called neon signs. It is relatively expensive because of its rarity. Argon In the late 19th Century, the composition of the atmosphere was supposedly well known. The components were oxygen, nitrogen , carbon dioxide, and varying amounts of water vapor. John William Strutt, Lord Reyleigh, was a scientist and professor at Cambridge. In firming up the densities of these gases, he found a puzzling inconsistency. Nitrogen prepared from ammonia had one density, but nitrogen from the air – after extracting the other gases – had a higher density. This discrepancy was tackled by the Scottish chemist William Ramsay. He found that the supposed nitrogen remnant from the atmosphere contained a portion that was not reactive. Its spectrum had unfamiliar lines; this was a new gas. The new element argon, "the lazy one", introduced a new column of the periodic table. Argon comprises about 1% of the atmosphere, which is more than carbon dioxide.

Its origin is rocks containing radioactive potassium that decays into argon. Argon is by far the most common noble gas, and therefore the cheapest. Krypton Krypton was discovered in 1898 as a trace element in atmospheric argon. William Ramsay and his assistant Morris William Travers detected unfamiliar spectral lines. The name krypton means "hidden". Xenon Xenon is the last gas discovered by Ramsay and Travers. From large quantities of argon, they extracted the heaviest fraction, krypton, and worked to separate it further. They found xenon in 1898, whose name means "stranger". Xenon is extracted from the air, where it occurs greatly diluted. It glows blue or violet in a gas discharge tube. As the heaviest safe inert element, it is used to fill premium lightbulbs and lamps that would otherwise be vacuum tubes, and it thereby slows the evaporation of the filament. Radon The heaviest of the noble gases is radioactive, a product of the decay of thorium and uranium . German physicist Friedrich Ernst Dorn in 1900 proved that radium minerals, themselves radioactive, emitted a radioactive gas, now called radon.

Uranium and thorium in the earth's crust are continually decaying into daughter products, one of which is the very heavy gas radon. Radon therefore tends to collect in certain low areas such as basements and especially mines. If it is inhaled, it presents a health risk, causing lung cancer. | Group Index | Main Index | Compounds | The Elements In 1785 the English scientist Henry Cavendish reported the results of some remarkably careful experiments. Just a few years after the discovery of oxygen, and before the concepts of atoms and molecules as we understand them today had been developed, he was able to show that there was a small component in air that was extremely unreactive. Although mains electricity was still a century away, scientists of this period did have primitive batteries and it was known that when an electric spark is passed through a mixture of nitrogen and oxygen they react to produce nitrogen dioxide. This brown gas is readily soluble in water to give a colourless solution of nitric and nitrogen monoxide, which reacts rapidly with oxygen itself producing more nitrogen dioxide so that ultimately all of the nitrogen dioxide is converted to nitric acid.

3NO2(g) + H2O(aq) � 2HNO3(aq) + NO(g) 2NO(g) + O2(g) � 2NO2(g) (It is this reaction taking place in car engines that is responsible for NOx pollution in the atmosphere and which makes a contribution to acid rain. These reactions are also used in the commercial manufacture of nitric acid.) Cavendish added more oxygen and repeated his reactions until he found that he had a small residue that did not react. We now know that this would have been argon with tiny traces of the other noble gases. What is a remarkable testament to his experimental skill is that his estimate of the quantity - 1 part in 120 - or 0.83% is very close to the true value of 0.93%. It was not until a century later that the problem was re-visited.

Another English scientist, Lord Rayleigh, a physicist interested in the properties of gases discovered, that the density of nitrogen extracted from the air by removing the oxygen and carbon dioxide from it was slightly higher than the density of nitrogen obtained by chemical means, such as by heating ammonium nitrite or alkali metal azides: NH4NO2(s) � N2(g) + 2H2O(l) 2NaN3 � 3N2(g) + 2Na(s) (The second of these two reactions is the one used to inflate car airbags - with the addition of additional reagents to mop up the molten sodium metal!). Rayleigh realised that his atmospheric nitrogen must be contaminated by another, heavier gas. In 1895 in an experiment akin to that of Cavendish, a colleague of Rayleigh's, William Ramsay, heated "atmospheric" nitrogen with magnesium, one of the few other elements that will react directly with nitrogen, and found that there was a small residue of a very unreactive gas.

3Mg(s) + N2(g) � Mg3N2(s) This residual gas was characterised by its physical properties such as boiling point, molecular mass and, most importantly, its distinctive atomic spectrum. Since all elements were known to have different atomic spectra it was this that proved that argon, as the new element was named, really was a new element and not just an unknown compound. The name argon is the Greek word for "lazy" and refers to its lack of chemical reactivity. It was at about this time that scientists were first able to achieve temperatures low enough to liquify air and by carefully separating the components of air on the basis of their different boiling points (fractional distillation) by 1898 Ramsay's research group had similarly characterised neon, krypton and xenon (named after the Greek words for "new", "hidden" and "strange" respectively.) Helium has a different history. It was first identified in the spectrum of the sun in 1868 and named after helios, the Greek word for the sun.

It wasn't known on earth until Ramsay identified it as a gas given out when some radioactive rocks are heated and by 1900 he had also obtained it from fractional distillation of air. When first named it was thought to be a metal, hence the ending "ium". The suggestion that it be renamed "helion" to match its true nature was eventually turned down on the grounds that "helium" was too well established, although this argument didn't prevent the change of the element symbol of argon from A to Ar in the 1950s! Noble Gas Compounds In 1933 Linus Pauling, one of the pioneers of modern bonding, reasoned that xenon with the lowest ionisation potential of the group was the most likely of the noble gases to form compounds and that if it was going to react with anything it would be with elemental fluorine, the most reactive of all of the elements. Following his suggestion, two American chemists, Yost and Kaye, mixed xenon and fluorine and passed an electric discharge through the mixture of gases.

They observed nothing, despite the fact that we now know that a reaction should take place under these conditions, and their failure set back noble gas chemistry by thirty years. In 1962, Neil Bartlett, a chemist at the University of British Columbia who was studying the chemistry of platinum hexafluoride, PtF6, had a lucky accident; his apparatus leaked letting in air. He noticed some unexpected yellow crystals and when they were analysed he found that the platinum hexafluoride was such a powerful oxidising agent that it was able to oxidise oxygen itself, forming the remarkable ionic compound O2+PtF6-. The hexafluoride had removed an electron from dioxygen, leaving the dioxygen molecular cation, O2+, which had never before been isolated in a stable compound. Shortly afterwards Bartlett noticed that the ionisation enthalpy for xenon was about the same as that for dioxygen; he tried mixing xenon gas and PtF6 and was rewarded by the instant formation of crystals which he formulated as Xe+PtF6-.

We now know that the crystals were more complicated than this and probably a mixture, the simplest compound of these three elements is XeF+PtF6-, which contains a diatomic cation with a similar electronic structure to that of a halogen molecule. At the same time as Bartlett's work Rudolph Hoppe, a German chemist, was reworking Pauling's ideas and concluded on thermodynamic grounds that xenon ought to form compounds with fluorine. Shortly after the publication of Bartlett's results he put his ideas to the test in the laboratory and prepared crystals of xenon tetrafluoride by passing an electric discharge through a mixture of the two elements. The results of Bartlett and Hoppe removed the mental block concerning the inertness of the noble gases and in a very few years the chemistry of xenon was found to be as extensive as that of many other elements. Why did Yost and Kaye fail? There are probably two explanations: mainly because the techniques and apparatus for handling these very reactive substances had improved enormously in the intervening years, but also because most chemists would have predicted that the product would have been a gas rather than a solid.

A few crystals could have been easily overlooked, especially since there were no transparent fluorine resistant materials available at that time. There is certainly no implication that they were any less than competent chemists. As we can see luck plays a part even in science. � H.J Clase 1999, 2001. Updated 01.01.15

Question: Which of the noble gases takes its name from the Greek for 'stranger'?
Answer: XENON